Concerted and Stepwise Proton-Coupled Electron Transfer for Tryptophan-Derivative Oxidation with Water as the Primary Proton Acceptor: Clarifying a Controversy

Concerted electron-proton transfer (CEPT) reactions avoid charged intermediates and may be energetically favorable for redox and radical-transfer reactions in natural and synthetic systems. Tryptophan (W) often partakes in radical-transfer chains in nature but has been proposed to only undergo sequential electron transfer followed by proton transfer when water is the primary proton acceptor. Nevertheless, our group has shown that oxidation of freely solvated tyrosine and W often exhibit weakly pH-dependent proton-coupled electron transfer (PCET) rate constants with moderate kinetic isotope effects (KIE ≈ 2–5), which could be associated with a CEPT mechanism. These results and conclusions have been questioned. Here, we present PCET rate constants for W derivatives with oxidized Ru- and Zn-porphyrin photosensitizers, extracted from laser flash-quench studies. Alternative quenching/photo-oxidation methods were used to avoid complications of previous studies, and both the amine and carboxylic acid groups of W were protected to make the indole the only deprotonable group. With a suitably tuned oxidant strength, we found an ET-limited reaction at pH < 4 and weakly pH-dependent rates at pH > ∼5 that are intrinsic to the PCET of the indole group with water (H2O) as the proton acceptor. The observed rate constants are up to more than 100 times higher than those measured for initial electron transfer, excluding the electron-first mechanism. Instead, the reaction can be attributed to CEPT. These conclusions are important for our view of CEPT in water and of PCET-mediated radical reactions with solvent-exposed tryptophan in natural systems.


I. Electrochemical Measurements
Tryptophan (W) radical formation is followed by rapid radical-radical dimerization, leading to self-inhibition which results in irreversible cyclic voltammograms (CVs). 1 Self-inhibition can be reduced by lowering the substrate concentration and increasing the scan rate. If self-inhibition is sufficiently minimized, reduction potentials can be determined from the variation of peak potentials with scan rates: 2 where E p is the peak potential, E° is the formal reduction potential, R, T, and F have the usual meaning, k dim is the rate constant of dimerization, C° is the concentration of substrate, and ν is the scan rate. Eq. S1 predicts a slope of 19.7 mV in a plot of E p vs. log(ν). The scan rate was varied between 0.1 and 5 V/s for a sample containing 0.2 mM the tryptophan analog WEE or NAWEE in 0.5 mM KP i and 0.1 M KNO 3 , Figure S1. This was the same range of scan rates used for a similar system. 2 At scan rates > 1 V/s the CVs started to lose their peaked shape, which was attributed to more capacitive current. We therefore decided to use scan rates from 0.1 V/s to 1 V/s. A linear fit to the apparent peak potentials in this range of scan rates yielded a slope 18 mV for both WEE and NAWEE, respectively, suggesting that self-inhibition was not significantly perturbing the CVs, Figure S2. 2 Figure S1. Cyclic voltammograms collected in 0.1 M KNO 3 and 0.5 mM KP i at pH 5.2 using a 2 mm glassy carbon working electrode, Ag/AgCl reference electrode, and a Pt counter electrode.

S3
Potentials are shown vs. Ag/AgCl. All peak potentials were measured vs. a Ag/AgCl reference electrode that was calibrated by measurement of the potential for the reversible [Fe(CN) 6 ] 3-/4couple (0.410 V vs. NHE). 3 A and C represent the same data collected with WEE with various zoom ins, B and D represent the same data collected with NAWEE but with different zoom ins. The black squares mark the peak potentials used in Figure S2. Conditions

ii. TA Kinetic Traces and Rate Constants at Low pH (< 4)
The reaction kinetics was studied as a function of WEE concentration probing at 450 nm where the consumption of [Ru(dmb) 3 ] 3+ is monitored by the [Ru(dmb) 3 ] 2+ bleach recovery. At low pH (< 4) the system is near to, or below, the pK a of oxidized tryptophan (≈ 4.3). Without deprotonation of the resulting radical, the electrochemical data in Table 1 of the main paper suggest that DG° » +100 meV for oxidation of WEE by [Ru(dmb) 3 ] 3+ , if the radical pK a -value is similar to that for W and NAW. We noticed that if a sufficiently large concentration of WEE was used (> 80 mM) the rate constants appeared single exponential ( Figure S7 and Table S2). Instead, at [WEE] = 50 mM and lower, the recovery did not exhibit single exponential kinetics, but showed biphasic behavior ( Figure S6). The fast rate component likely represents a single ET pre-equilibrium reaction, where the pre-equilibrium position is shifted depending on the concentration of reactants. The slow component can be attributed to follow-up reactions such as PT and radical dimerization, that drive the reaction to completion. The relative amplitude of the fast (single exponential) component decreased with decreasing [WEE], and was very small for [WEE] = 5 mM (Table S1).
At [WEE] = 25 and 50 mM, the biphasic kinetics observed was well fitted by a double exponential decay. At [WEE] = 5 mM, the slow phase of the bleach recovery was fitted with a term that was second order in [WEE • ]. The rate constant for the fast component had to be locked because of its small amplitude, and a value for the pseudo-first order rate constant was used that was calculated from an average of the second-order rate constant obtained from a fit to the 25 and 50 mM rate constants. We attempted to fit the 25 and 50 mM kinetic traces to single exponential plus second order recovery, however this did not converge properly, likely because the amplitude of the second phase is relatively small.
Even at the lowest concentrations of WEE employed in this pH region (5 mM), its concentration is significantly larger than the concentration of Ru II on the product side ([Ru II ] ≈ 25 µM), so the rapid phase can be assigned to setting the equilibrium (Eq. S2), with an observed rate constant equal to the sum of the forward and backward ET rate constants. The slower component of the fits represents the further reaction in Eq. S3 that consumes both Ru III and WEE •+ . At [WEE] ³ 80 mM, the equilibrium in Eq. S2 is shifted sufficiently to the right that Ru III is essentially consumed already in the first phase, and the observed rate constant is close to that for the forward ET step. With [WEE] = 25 mM, about half of the Ru III is consumed in each phase, seen by the similar relative amplitudes in Table S1.
At these low pH values, the protonated tryptophan radical (WEE •+ ) can be observed ( Figure S7). Its absorption is centered at 560 nm, but it also absorbs (slightly less) at 510 nm. A comparison in S7 TA amplitude between the two wavelengths reveals that primarily the protonated radical is formed on the timescale of the experiment, Figure S7. 7 Table S1. [WEE] (mM) iii. TA Kinetic Traces and Rate Constants at Intermediate pH (5 < pH < 7) Figure S8 show the transient absorption (TA) single shot kinetic traces with single exponential fits (orange) used to extract the rate constants seen in Figure 5, main text.

iv. TA Kinetic Traces and Rate Constants at High pH (> 7)
Excitation of [Ru(dmb) 3 ] 2+ with a 355 nm 50 mJ laser pulse results in two-photon ionization which produces Ru(III) and solvated electrons (esolv ). There are several species that can absorb at the wavelengths where we typically monitor the WEE oxidation (450 nm for Ru(II) bleach and 510 nm for WEE • formation), extinction coefficients and reaction steps with difference extinction coefficients for each reaction, and rate constants are found in Table S4 and Scheme S1, respectively. The esolv exhibits a broad absorption band centered around 700 nm that extends towards the blue part of the spectrum. 9 The scheme does not include reaction between esolv and WEE • because the solvated electrons were found to decay before WEE • was formed (vide infra).  (3) is expected to dominate. This is indeed what is observed in Figure S9, where the electron signal (broad absorption around 700 nm), decays on the time scale of 200 -1000 µs as Ru(II)Lis formed (band maximum around 510 nm). The 450 nm Ru(II) bleach recovery shows a biphasic behavior, with a strong ~300 ns component that is followed by a much slower recovery on the time scale of 10 µs ( Figure S15). The fast component is due to decay of the *Ru(II) that did not absorb a second photon, but the remaining bleach in the slow component shows that ca. 1 µM Ru(III) is formed. Kinetic traces recorded at 450 nm and 540 nm with only [Ru(dmb) 3 ] 2+ in 0.5 mM KP i indicate that the 450 nm bleach recovery occurs with the same rate constant as for the Ru(II)Ldecay, k ≈ 7´10 3 s -1 from a fit to a first order recovery at 540 nm and 450 nm, Figure S10. This indicates that without WEE in solution, Ru(II)Lmainly decay via recombination with Ru(III). One would expect the decay to exhibit second order kinetics, however such a fit did not converge, probably because reaction (4) is not completely dominating. The first order fits presented below are not perfect, but can be used to indicate similarities or differences between decays at different wavelengths and pHvalues. When 0.1 mM WEE is added to the solution, the 450 nm recovery is accelerated and the decays fit a single exponential.

(left) and pH = 9.3 (right). The orange line represents a single exponential fit, with the fit residuals shown in the bottom graph. The traces show that recombination between Ru(III) and Ru(II)Ldoes not exhibit any pHdependence.
In the presence of WEE, the 450 nm recovery is accelerated, and exhibits first order kinetics with k = 1.4´10 4 s -1 at pH = 7.6; Figure S12 and Table S5 show traces and data for all pH values studied using this method. The rate constant should represent the sum of the rate constants for reactions (2) and (4). It is clear that the 450 nm recovery becomes faster as the pH increases. Subtracting the rate constant for 450 nm recovery without WEE (reaction (4) in Scheme 1) from the rate constants with WEE gives a corrected rate constant (k corr ) which represents reaction (2) in Scheme S1. These are the rate constants found in Figure 5 in the main paper. The rate of 450 nm bleach recovery as a function of pH in the absence of WEE was also determined. The kinetic traces, seen in Figure S11, show that this recovery is independent of pH, showing that the pH-dependent rate constants are due to reaction with WEE.    Figure 5, main text.
When WEE is present in the solution, the observed rate constants determined at 510 nm (which are dominated by Ru(II)Ldecay) is obtained as ~1.5´10 4 s -1 ( Figure S13 right panel), and the kinetics is independent of pH. This is slower than the disappearance of Ru(III), which shows that another oxidized species is formed to balance the electrons on Ru(II)L -, which is attributed to WEE • . Some absorption around 450 nm from WEE • can also explain why the 450 nm traces return to baseline, ( Figure S13, left panel) even though Ru(II)Lshows a bleach at this wavelength. The To ensure that the 450 nm recovery was first order with respect to [WEE], the rate using 1 mM WEE at pH 8.8 was measured, Figure S14. The bleach recovery exhibited double exponential recovery with k 1 = 3.0´10 6 s -1 (*Ru(II) decay) and k 2 = 1. no WEE S19
[Ru(dmb) 3 ] 3+ was generated with [Co(NH 3 ) 5 Cl] 2+ as an irreversible quencher. The rate constants of NAWEE oxidation exhibited no dependence on pH and are in good agreement with previously published data for NAWEE using the same photosensitizer but with [Ru(NH 3 ) 6 ] 3+ as the quencher, 6 as seen in Figure S16. The pH independent rate constants determined with [Ru(dmb) 3 ] 3+ and [Ru(NH 3 ) 6 ] 3+ as the photosensitizer and quencher, respectively, are indicative of an ET limited reaction, i.e. an ETPT mechanism.  [ZnTPPS] 4was excited in the Q-band (545 nm) and spectral changes were studied at 470 nm, in the absorption band of the oxidized form of the porphyrin. The kinetic traces recorded at 470 nm in the presence of NAWEE fit to a double exponential function that did not return to the baseline.

S21
Upon laser flash-photolysis, the steady state UV-Vis spectrum of [ZnTPPS] 4changed, showing increased intensity between 440 -470 nm, Figure S17. This indicated that [ZnTPPS] 4degrades during the TA experiment, the products of which absorbed at 470 nm. For this reason, all kinetic traces used are from the first laser shot to a fresh sample. The difference spectra in the inset of Figure S17 show the changes in the UV-Vis spectra after degradation.  Table S6. Table S6. Rate constants, pre-exponential factors and baseline offset (Y 0 ) obtained from a double exponential fit to TA traces measured at different probe wavelengths, seen in Figure S19.  Figure S20 and Table S7 at pH 6.3 -9.5. Using a weaker oxidant led to slower kinetics, and NAWEE had limited solubility in water that hindered us from collecting faster observed pseudo-first order rate constants. Therefore, kinetic traces for this system were collected on very long timescales up to 10 s. Extra precautions were taken to accommodate collection of TA traces on long timescales. Signals settings were optimized using a dummy sample so that each kinetic trace was collected on a fresh sample that had not been exposed to light. Exposure to excess probe light was limited during detection by filtering the probe light through a monochromator such that FWHM was 18.6 nm before the sample. Taking this step ensures minimal initiation of the photosensitizer/quencher reaction which accelerate [ZnTTPS] 4degradation. Xe lamps fluctuate in intensity on longer timescales. This unavoidable effect can be seen in traces collected on longer timescales in Figure S20. To minimize random fluctuations of the Xe lamp, the temperature was equilibrated for at least 30 minutes up prior to measurements. Importantly, the fluctuations are random and good exponential fits are still obtained.  Table S7.

iv. Rate Constants for NAWEE Oxidation as a Function of Buffer Concentration
TA traces have been measured in samples containing different concentrations of buffer from 0.5 to 5 mM. Rate constants have been calculated from the short component of a double exponential fit to assess the contribution of the buffer as an alternative proton acceptor in the PCET reaction. The results are shown in Table S8. One order of magnitude increase in buffer concentration corresponds to less than 20% increase in rate constants, and demonstrates that at 0.5 mM concentration, the role of the buffer as proton acceptor is negligible.

i. Concerted PCET
For OHto act as a proton acceptor, it must be present at sufficient concentration to account for the observed rate constants. Stepwise mechanisms can be excluded; PTET is excluded on the basis of the high pK a value of the indole proton (~17), and ETPT is excluded on the basis of the rate constants that increase above the value obtained at pH < 4.5 where the mechanism is limited by ET. We therefore only consider the concerted mechanism. For CEPT, we can analyze the reaction as OHreacting with the {PS ox ⋯ W} encounter complex (PS ox = oxidized photosensitizer) after it has formed, see below.
k OH-in reaction S5 has a sufficient driving force that it can be treated as an irreversible step, albeit slow compared to the pre-equilibrium in S4 because of the low [OH -]. By applying a steady state approximation on the encounter complex, we get the following expression for the rate: where the pseudo first order rate constant k' obs is defined as: Assuming that there is no specific driving force for formation of the complex in reaction S4, i.e., ΔG° = 0 we get k d /k -d = 1 M -1 . Given that k d ≈ 10 10 M -1 s -1 in water, this allows us to estimate k -d to ≈ 10 10 s -1 . In solution, k OH-cannot be greater than the rate of diffusion, k diff ≈ 10 10 M -1 s -1 in water, and is in fact likely lower. Comparing k -d to k OH-[OH -] would even at pH = 10 yield k OH-[OH -] < 10 6 s -1 << k -d . This assumption allows us to simplify Eq. S7 to Eq. S8, below.
Recalling that k d /k -d = 1 M -1 , Eq. S8 can be further simplified to k' obs = k OH-[OH -]/M (k' obs is a second-order rate constant) which makes comparison with the experimentally observed k PCET possible and should give k PCET = k OH-[OH -] if OHis the primary proton acceptor. As k OH-is ≤ k diff (≈ 10 10 M -1 s -1 ), [OH -] is not sufficiently large even at pH = 11.4 to yield rate constants that are similar to what was experimentally observed (k PCET = 3.5´10 8 M -1 s -1 at pH = 11.4) with WEE and

S28
[Ru(dmb) 3 ] 3+ as oxidant. For NAWEE, the same is found for pH ³ 9. Thus, we conclude that OHcannot be the primary proton acceptor in the pH range examined.
ii. Stepwise PTET For a PT limited PTET reaction the derivation of k PTET is analogous to the one for k' obs . The rate constant for PT decreases by one order of magnitude for each unit of ΔpK a . The deprotonation of W with OHas proton acceptor have a ΔpK a = -3. k PTET can therefore be maximum 10 7 M -1 s -1 , which is not enough to account for the observed rate constants for neither WEE nor NAWEE. With the same logic, PT limited PTET with water as proton acceptor (ΔpK a = -17) can also be ruled out.